Introduction to Electrochemistry Read Chapter 14. Electrochemistry can be broadly defined as the study of charge-transfer phenomena. As such, the field of electrochemistry includes a wide range of different chemical and physical phenomena. These areas include (but are not limited to): battery chemistry, photosynthesis, ion-selective electrodes, coulometry, and many biochemical processes. Although wide ranging, electrochemistry has found many practical applications in analytical measurements. Electro-analytical Chemistry Electro-analytical chemistry is the field of electrochemistry that utilizes the relationship between chemical phenomena which involve charge transfer (e.g. redox reactions, ion separation, etc.) and the electrical properties that accompany these phenomena for some analytical determination. This relationship is further broken down into fields based on the type of measurement that is made. Potentiometry involves the measurement of potential for quantitative analysis, and electrolytic electrochemical phenomena involve the application of a potential or current to drive a chemical phenomenon, resulting in some measurable signal which may be used in an analytical determination. There are two parts to understanding electrochemistry: the first is thermodynamics; the second is the kinetics of electrode processes. For the latter, one needs to study the surface chemistry to obtain a real understanding of how electrochemical systems work. For chemists to understand the principles underlying the functioning of such practically important systems as batteries, fuel cells, corrosion, electrolysis, as well as membranes and biomembranes (of utmost importance for the understanding of, drug delivery and the functioning of cell membranes) they must therefore be taught the basics of interfacial structure, electrochemical kinetics and transport processes. (From International Society of Electrochemistry) site Example of a Galvanic Cell Redox reactions involve electron transfer. Acid-Base reactions involve proton transfer. The key difference is that electrons can be transported through space via wires or any other conducting device with the result that the oxidation and reduction reactions can occur in different places. Protons are transported in an aqueous (or some other polar) environment, so acid-base reactions occur in the same place. Basic Concepts Redox reactions involve a species which is oxidized and another that is reduced. Fe3 V 2 Fe2 V 3 In the above, Fe3+ is reduced to Fe2+ . It is the oxidizing agent. Since DG<0 for this reaction we can say that V3+ wants the extra electron less than Fe3+. V 3 Fe3 2 2 Galvanic Cell Cd 2Ag An Aside Why won’t this cell work? Ag+ will go to left electrode and ask for e from Cd(s) directly. Will this cell work? How badly to the electrons want to flow? q I E/R sec I = current in amps R = resistance in ohms E = potential difference in q = n x F Volts Coulombs Moles Coulomb DG -nFE Moles Voltaic Cells Electrochemical cells that use an oxidation-reduction reaction to generate an electric current are known as galvanic or voltaic cells. Let's take another look at the voltaic cell in the figure below. Voltaic Cells Taken from http://chemed.chem.purdue.edu/genchem Within each half-cell, reaction occurs on the surface of the metal electrode. At the zinc electrode, zinc atoms are oxidized to form Zn2+ ions, which go into solution. The electrons liberated in this reaction flow through the zinc metal until they reach the wire that connects the zinc electrode to the platinum wire. They then flow through the platinum wire, where they eventually reduce an H+ ion in the neighboring solution to a hydrogen atom, which combines with another hydrogen atom to form an H2 molecule. The electrode at which oxidation takes place in a electrochemical cell is called the anode. The electrode at which reduction occurs is called the cathode. The identity of the cathode and anode can be remembered by recognizing that positive ions, or cations, flow toward the cathode, while negative ions, or anions, flow toward the anode. In the voltaic cell shown above, H+ ions flow toward the cathode, where they are reduced to H2 gas. On the other side of the cell, Cl- ions are released from the salt bridge and flow toward the anode, where the zinc metal is oxidized. The voltaic cell consist of the two reactions. Zn(s) Zn 2 2e oxidation + 2H 2e - H (g) reduction 2 2H Zn(s) Zn 2 H 2 (g) Or equivalently we can write the reactions as follows 2H 2e - H 2 (g) Zn 2 2e Zn(s) 2H Zn(s) Zn 2 H 2 (g) We can only measure E for the full reaction. We would like to calculate E for the half reactions. Before doing this, we must recognize the E depends on concentrations. Voltaic Cells Zn 2 1 H 1 H2 (g) 1 Taken from http://chemed.chem.purdue.edu/genchem Since reactants and products are in their standard states, we call the E for this cell the standard reduction potential (Eo). Here Eo = .76V. 2 Zn(s) Zn 2e + 2H 2e - H 2 (g) 2H Zn(s) Zn 2 H 2 (g) We arbitrarily define the potential for, one half reaction, the second reaction above to be exactly 0V when reactants and products are in their standard states. Since Eo for the cell is the sum of Eo’s for the two half reactions we see that Eo for the first half reaction is .76V. Oxidizing Power Increases Voltaic Cells Zn 2 1 H 1 H2 (g) 1 Taken from http://chemed.chem.purdue.edu/genchem Zn(s) Zn 2 2e + 2H 2e - H 2 (g) 2 2H Zn(s) Zn H 2 (g) This voltaic cell on the previous slide is fully described with the following notation Zn(s) | Zn 2 (aq, Α 1) || H (aq, Α 1) | H 2 (g, Α 1) | Pt(s) Zn(s) | Zn 2 (aq, Α 1) || S.H.E. Line Notation For Voltaic Cells Voltaic cells can be described by a line notation based on the following conventions. Single vertical line indicates change in state or phase. Within a half-cell, the reactants are listed before the products. Activities of aqueous sol’ns are written in parentheses after the symbol for the ion or molecule. A double vertical line indicates a junction between half- cells. The line notation for the anode (oxidation) is written before the line notation for the cathode (reduction). The line notation for a standard-state Daniell cell is written as follows. Electrons flow from the anode to the cathode in a voltaic cell. (They flow from the electrode at which they are given off to the electrode at which they are consumed.) Reading from left to right, this line notation therefore corresponds to the direction in which electrons flow. Zn | Zn2+(1.0 M) || Cu2+(1.0 M) | Cu anode cathode (oxidation) (reduction) The Nernst Equation (14-4 of book) The Nernst equation relates the potential of a cell in its standard state to that of a cell not in its standard state. Consider the reaction below. Zn 2 2e Zn(s) 2 2Zn 4e 2Zn(s) We know from Le Chatelier’s principle that increasing the concentration of Zn2+ should drive the reaction to the right. In other words it should decrease the potential of the half cell. The Nernst equation allows us to calculate this increase for the above half reaction as RT Zn(s) RT 2 EE o ln EE o ln Zn(s) 2F Zn 2 4F 2 Zn 2 The Nernst Equation Continued The Nernst Eq. for the reaction bB ne cC is RT c EE o ln C nF b B At 25oC this equation simplifies to .05916V c EE o log C n b B The Nernst Equation For Complete Cell E E E Here E+ and E- are the potentials of the half cells connected to the positive and negative terminals of potentiometer respectively. Let’s consider an example. Voltaic Cells - + Zn 2 1 H 1 H2 (g) 1 Taken from http://chemed.chem.purdue.edu/genchem E E E E+ and E- are potentials of half cells connected to positive and negative terminals of potentiometer respectively 2H 2e - H 2 (g) Zn 2 2e Zn(s) 2H Zn(s) Zn 2 H 2 (g) .05916V H2 ( g ) .05916V Zn(s) E log 2 - - .76 log 2 H 2 Zn 2 .05916 H 2 ( g ) Zn 2 See page 295 for E .76 log another example 2 2 H Calculating Equilibrium Constants Consider the reaction Ce 4 Fe2 Ce3 Fe3 made up of the following two half reactions Ce 4 e Ce3 Eo=1.700V 3 2 Fe e Fe Eo=0.767V Since Eo is greater for cerium this reaction will be the reduction reaction. The standard potential for the galvanic cell would be E E - E 1.700 - (0.767) o o o Calculating Equilibrium Constants Continued Consider the reaction Ce 4 Fe2 Ce3 Fe3 In a galvanic cell we would have 3 3 RT [Fe ][Ce ] E E - E- E - E - 0 0 ln nF [Fe 2 ][Ce 4 ] At equilibrium E=0 and RT [Fe 3 ][Ce 3 ] Eo - Eo ln 2 4 .05916 log K at 25 o C F [Fe ][Ce ] K 1016 DG o Eo - Eo This connection to free nF energy is important Calculating Equilibrium Constants for Nonredox Reactions (14-5) 2 Consider the reaction FeCO 3 ( s ) Fe CO 3 2- This is a Ksp problem. Not a redox problem. Nonetheless we can use electrochemistry to calculate Ksp by considering FeCO3 ( s) 2e Fe(s) CO 2- 3 E 0.756V o Fe2 2e Fe(s) E o 0.44V E -E o o .059 2 2 log [Fe ][CO3 ] 2- .059 2 log K sp 2- .316 2.316 / .059 11 log K sp Ksp 10 10 .059 (at 25oC) Latimer Diagrams (Box 14-2 of book) The oxidation states of elements are related to each other IO IO HOI I 2 I - 4 1.589 - 3 1.154 1.430 .535 - (+7) (+5) 1.299 (+1) (0) (-1) 2H IO 2e IO H 2 O - 4 - 3 DG o 1 5H IO 3 4e HOI 2H 2 O - DG o 2 7H IO -4 6e HOI 3H 2O DG o 3 2 1.589 4 1.154 DG DG DG o 1 o 2 o 3 1.299 6 2E 1 4E o o n1FE1 n 2 FE 2 n 3FE 3 o o o 2 E3 o 6 Electrochemistry Skills • Understand how voltaic cells work. • Be able to calculate standard reduction potentials for voltaic cells, given the chemical reactions. • Be able to describe a voltaic cell using the line notation and visa versa. Know which way electrons flow and where the anode and cathode are. • Know how to work with the Nernst Eq. to include concentration dependencies and calculate equilibirum constants. • Know the relation between E and DG and can use this relation in constructing Latimer diagrams.
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